CHAPTER 11  CHEMICAL QUANTITIES


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1 I. THE MOLE CONCEPT CHAPTER 11  CHEMICAL QUANTITIES A. What is a mole? 1. a mole is the SI unit of measurement of counting; just like a dozen is a measurement 2. a mole is a number a x is called Avogadro's number b. 1 mol = 6.02 x representative particles rep. particle is a generic term. There are three specific terms: Atom If the particle in the problem is an element then the specific term is an atom. For example: Molecule If the particle in the problem is a covalent compound, i.e. starts with a nonmetal, then the specific term is molecule. For example: Formula Unit If the particle in the problem is an ionic compound, i.e. starts with a metal or ion, then the specific term is formula unit. For example: B. Converting moles to particles and particles to moles 1. Remember any conversion factor can be written two ways: So using Avogadro's number we can write: 2. to solve a problem follow the same steps for dimensional analysis: Do you remember them? Write them below. Example 1 We will be dealing with CO. So will the units for Avogadro's number be atom, molecule or formula unit? How many molecules of CO are in 3.0 moles of CO. Example 2 We will be dealing with Zn So will the units for Avogadro's number be atom, molecule or formula unit? Calculate the number of moles that contain 4.50 x atoms of Zn? 1
2 II. Mass and the Mole A. Molar mass  The mass in grams of one mole of any pure substance. B. Molar Mass of an element is equal to its atomic mass. For example, the molar mass of oxygen is equal to C. Round all atomic masses from the periodic table to 3 significant figures. D. The unit for molar mass is gram/mole E. Mole to Mass Conversion and Mass to Mole Conversion This time our conversion unit will be: Example1: If you have 2.0 moles of lithium, how many grams do you have? Example 2: If you have 205 grams of carbon, how many moles do you have? III. Moles of Compounds A. Chemical Formulas and the Mole Chemical formula for a compound indicates the types of atoms and the number of each contained in one unit of the compound. For example: Freon, CCl 2 F 2 In one mole of freon, you would have moles of carbon, moles of chlorine, and moles of fluorine. We will need to convert from moles of a compound to moles of individual atoms in the compound or from moles of individual atoms in a compound to moles of the compound. The following conversion factors can be written for use in converting. 2
3 For instance: How many moles of fluorine atoms are in 5.50 moles of Freon. Example 1: Determine the number of moles of Al 3+ in 1.25 moles of aluminum oxide. Example 2: How many moles of oxygen atoms are in 2.5 moles of diphosphorous pentoxide. B. Molar Mass of Compounds 1. the molar mass of a compound can be determined by first finding the mass of each element in the compound from the periodic table. The mass of each element is then multiplied by the number of atoms of that element in the chemical formula and then the products are added together. Example #1: molar mass of hydrogen (diatomic molecule  BrINClHOF) Example #2: molar mass of H2O Example 3: Calculate the molar mass of calcium hydroxide. C. MoleMass conversions and Mass to mole conversions Our conversion unit will be: Example #1. Calculate the mass of 2.50 moles of allyl sulfide, (C 3 H 5 ) 2 S 3
4 Example #2. Calculate the moles in 922 g of iron (III) oxide. D. Mass to Particles Conversion and Particles to Mass When no mole is mentioned in the problem, both conversion factors will be required. We call it the chemistry twostep. Example 1: How many formula units are in 55.2 g of LiCl? Example 2: What is the mass in grams of each of 1.00 x atoms of Mn? SECTION 11.4 EMPIRICAL AND MOLECULAR FORMULAS I. Percent Composition A. Defined: the percent by mass of each element in the compound B. Determining the percent composition of a known compound 1. determine the molar mass of each element in the compound and the molar mass of the entire compound 2. divide the molar mass of each element by the molar mass of the entire compound and then multiply by 100. Sample Problems: Calculate the % composition of 4
5 a. C 2 H 6 b. calcium acetate: Ca(C 2 H 3 O 2 ) 2 You try! II. Empirical and Molecular Formulas 1. Empirical Definition: formula giving the lowest wholenumber ratio of the elements in a compound 2. may or may not = molecular formula (actual # of atoms) 3. example: molecular formula for dinitrogen tetrahydride = N 2 H 4 empirical formula = NH 2 molecular formula for carbon dioxide = CO 2 empirical formula = CO 2 4. Molecular Formula Definition: actual # of atoms in a molecule 5. Steps to calculating empirical and molecular formulas a. Write each element down in column form. b. If given %, assume 100 grams This means that 25% is 25 grams. So instead of %, write the unit g. c. Divide the number of grams of each element by its molar mass to determine moles. d. Choose the least moles and divide all of the moles by it. e. If you get whole numbers, then these are the subscripts for each element in the compound. If you don't get whole numbers, then multiply all by 2, 3, or 4 to get a whole number. f. Write the empirical formula, each element with its subscript. g. To determine the molecular formula, calculate the molar mass of the empirical formula. Compare it to the molar mass of the compound which must be given. If they are equal then the empirical formula and the molecular formula are the same. If not, divide the larger by the smaller. You should get a whole number. Multiply the subscripts of the empirical formula by the whole number and rewrite the formula with the new subscripts. 5
6 Sample problems 1. What is the empirical formula for a compound that contains 10.89% magnesium, 31.77% chlorine, and 57.34% oxygen? 2. A compound with a formula mass of amu is found to be 85.64% carbon and 14.36% hydrogen by mass. Find its molecular formula. You try! 6
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